A buffer solution is a type of solution able to resist a change in pH due to the addition of a little amount of acid or alkali or due to dilution.
How buffers work
A buffer solution is a mix of acid and alkali: the acid neutralises the alkali and the alkali neutralises the acid. However, a strong acid and a strong alkali cannot be used as they would react with one another thereby not acting as an effective buffer. Neither can a very weak acid and a very weak alkali as they will not react enough.
Instead, acids and bases should be chosen that are strong enough to react with H3O+ and OH– ions but not with one another. A weak acid and its conjugate base are an ideal mixture, for example NH4Cl and NH3.
Resisting pH change: adding acid and alkali
Taking the example above, this mixture will react with an acid and an alkali as follows:
- NH4+(aq) + OH–(aq) ? NH3(aq) + H2O(l)
- NH3(aq) + H3O+(aq) ? NH4+(aq) + H2O(l)
Resisting pH change: dilution
With dilution, the weak acid and weak base are both able to dissociate more in order to compensate:
- NH4+(aq) + H2O(l) ? NH3(aq) + H3O+(aq)
- NH3(aq) + H2O(l) ? NH4+(aq) + OH–(aq)
Other buffer solutions
A buffer a mixture does not have to be a weak acid and its conjugate base. Instead, you can use any mixture of a weak acid and a weak base.
A substance which is able to act as a weak acid and a weak base can also behave like a buffer. For example sodium hydrogencarbonate:
- HCO3–(aq) + H3O+(aq) ? CO2(g) + 2H2O(l)
- HCO3–(aq) + OH–(aq) ? CO32-(aq) + H2O(l)
Amino acids can also act like buffers, 2-aminopropanoic acid for example:
- CH3CH(NH2)COOH(aq) + OH–+(aq) ? CH3CH(NH2)COO–(aq) + H2O(l)
- CH3CH(NH2)COOH(aq) + H3O+(aq) ? CH3CH(NH3+)COOH(aq) + H2O(l)
Calculating the pH of buffer solutions
If the buffer solution is composed of a weak acid and its conjugate base then it is possible to calculate its pH using the following method.
Take the mixture: CH3COOH / CH3COONa
Therefore: Ka = ([CH3COO–][H3O+]) / [CH3COOH]
So: [H3O+] = Ka[CH3COOH] / [CH3COO–]
In general it is possible to calculate the concentration of H3O+ ions from:
[H3O+] = Ka[acid] / [base]
To create a buffer solution of a specific pH the concentrations of the acid and base relate to a particular ratio. This can be worked out by expressing the pH in terms of the concentrations of the reactants:
[H3O+] = Ka[acid] / [base]
So: log10[H3O+] = log10 Ka + log10([acid] / [base])
So: – log10[H3O+] = -log10 Ka + log10([base] / [acid])
So: pH = pKa + log10([base] / [acid])
Limiting changes in pH with a buffer solution
It is possible to quantitatively show how a buffer solution is able to limit changes in pH.
Take the following mixture: 0.6M HClO (Ka = 3.7 x 10-8 M) and 0.2M NaClO with a pH of 7.0
When 0.01 moles of HCl are added to 100cm3 of the buffer:
|
ClO–(aq) |
+ |
H3O+(aq) |
? |
HClO(aq) |
+ |
H2O(l) |
| Initially |
0.02 |
0.01 |
0.06 |
|||
| Finally |
0.01 |
– |
0.07 |
Therefore: the [base] / [acid] ratio = 0.01 / 0.07
pH = pKa + log10([base] / [acid]) = 6.58 (a unit change of 0.4)
If 0.01 moles of HCl were added to 100cm3 of pure water then the resulting solution would have a pH of 1.0, a change of 6 units.
When 0.01 moles of NaOH are added to 100cm3 of the buffer:
|
HClO(aq) |
+ |
OH–(aq) |
? |
ClO–(aq) |
+ |
H2O(l) |
| Initially |
0.06 |
0.01 |
0.02 |
|||
| Finally |
0.05 |
– |
0.03 |
Therefore: the [base] / [acid] ratio = 0.03 / 0.05
pH = pKa + log10([base] / [acid]) = 7.21 (a unit change of 0.2)
If 0.01 moles of NaOH were added to 100cm3 of pure water then the resulting solution would have a pH of 13.0.
Therefore, when a small amount of acid or alkali is added to a buffer solution the pH does not change by much.
However, if too much of an acid or an alkali is added then the buffering capacity of a buffer solution can be exceeded leading to a massive pH change. In the example above this would be over 0.02 moles of HCl or over 0.06 moles of NaOH.
A buffer solution is most effective at resisting a pH change if the concentrations of acid and alkali are equal. If there is more acid than base or more base than acid in a buffer solution then it cannot resist change on the addition of an acid or base respectively.
Natural buffers
A number of biological systems reply on a relatively constant pH:
- Blood: the pH must be kept around 7.4 for which hydrocarbonate ions are used to maintain this level.
- Tears: the pH must also be kept around 7.4 and amino acids are used to maintain this level.